General Chemistry

1. General Chemistry • Assistant professor • Mervat Mohamed Hosny

2. 6-Quantum mechanical atom (Schroedinger)

3. 1-Democritus: • He theorized that all matter is composed of • small indivisible particles called atoms

4. :2-Dalton’s atomic theory • *each element is composed of minute • indivisible particles called atoms • *all atoms of a given element are chemically identical to each other ,atoms of one element are different from the atoms of all other element

5. 3- during ordinary chemical reactions atoms of one element cannot be changed into atoms of different element. • 4- atoms are not created or destroyed • 5- compound is formed when atoms of more than one element combine

6. 3-J.J.Thomson-CRT • *he discovered the electron • * in thomson‘s model ,electrons are • embedded in a positive sphere of matter

7. 4-Rutherford gold foil experiment: • *he established that the positive charged alpha particles emitted by certain radioactive elements (helium) . • * he used these alpha particles to establish the nuclear nature of atoms . • *in these experiments ,he directed a stream of positive charged helium ions (alpha particles)at a very thin sheet of gold foil

8. conclusion • *most of the mass and all of the positive charge of the atom are contained in a small • space called the nucleus • *most of the volume of the atom is empty space occupied by tiny negatively charged electrons

9. *negative charged electrons outside nucleus =positive charge inside nucleus • *the atom is electrically neutral • *protons:+vely charged subatomic particles found in nucleus • *neutrons : neutral (uncharged)subatomic particles found in nucleui • *electrons very small – vely charged subatomic particles

10. 5- The Bohr model: • *electrons in an atom exist in specific regions at various distances from the nucleus . • *The electrons are rotating in orbits around the nucleus like planets rotating around the sun.

11. *he describe hydrogen atom as a single electron rotating in an orbit about a relatively • heavy nucleus . • *he applied the concept of energy quanta, proposed by the German physicist Planck

12. Planck stated that : • *energy is never emitted in a continuous stream but only in small discrete packets • called quanta

13. Bohr theorized that : • *There are several possible orbits for electrons at different distances from the nucleus • *but electron had to be in one specific orbit or another .

14. *It could not exist between orbits • *when a hydrogen atom adsorbed one or more quanta of energy ,its electron jumped to another orbit a greater distance from the nucleus. • *when the electron fell back to lower orbits ,it emitted quanta of energy as light ,giving rise to the spectrum of hydrogen .

15. *each orbit is at a different energy level • *an electron in the orbit closest to the nucleus • is in the 1st energy level ,at greater distances it • may be in the second ,3rd or fourth energy level

16. 6-Quantum mechanical atomSchroedinger) ) • *They found that Bohr’s assumptions have to be modified • *Difficulty arise in applying the theory to atoms containing many electrons • *Bohr’s concept was replaced by quantummechanics theory

17. One of the chief difference between the 2 theories is that: • In the quantum mechanics theory electrons • are not considered to be revolving around the • nucleus in orbits but to occupy orbitals cloud • like regions surrounding the nucleus and • corresponding to energy levels

19. Erwin Schrodinger introduced his famous wave equation Quantum mechanics or wave mechanics • He describe an electron as simultaneously having properties of: • 1-a wave (like light) • 2-and a particle (have mass)

20. The solution of the Schrodinger equation is complex but as aconclusion: • There is four quantum numbers which • define the location and properties of • electrons in atoms: n,l,m,s

21. n is the principle quantum no indicate the energy levels of the electron relative to their distance from the nucleus • n=1,2,3,……… • But always 1-7 • بعد الالكترون عن النواة

22. L=2nd quantum no explain the shape oforbital • Electron exist in orbitals having specific shapes • S P d f

23. 3-m magnetic quantum no • Orientation in space • *electron orbitals have specific orientation in • Space • *This quantum number accounts for the number of s,p,d,f orbitals that can be present in the principal energy level

24. 4-Spine quantum no (s) • *an electron spins about its own axis in either a clockwise or counter clockwise direction • *S relates to the direction of spin of an electron • *when 2 electrons occupy the same orbital,they must have opposite spins

25. *when an orbital contain 2 electronsthe electrons are said to be paired • NO ELECTRONS IN AN ATOM CAN HAVE • THE SAME 4 QUANTUM NUMBERS

26. 7- Energy levels of electrons • *all the electrons in an atom are not located the same distance from the nucleus • *as said in Bohr theory and quantum mechanics the probability of finding the electrons is greatest at certain specified distance s called energy levels, from the nucleus

27. *energy levels are also referred to as electron shell and may contain only a limited number of electrons • * energy levels are numbered startly with n=1 to n=7 • Or K,L,M,N,O,P,Q • Where K=1st energy level,L=2nd energy level

28. *the maximum number of electrons that can occupy a (specific energy level =2(nxn • n= number of the principle energy level • E.g. for shell k or energy level 1=2x(1x1)=2 • E.g. for shell l or energy level 2 =2x(2x2)=8

29. 8- energy sublevels of electrons: • *the principle energy levels contain sublevels designated by the letters s,p,d,f • *s sublevels consists of 1 orbital • *p sublevels consists of 3 orbitals • * d sublevels consists of 5 orbitals • * f sublevels consists of 7 orbitals

30. The maximum no of electrons that can exist in these sublevels is : • S sublevel 2 electrons • P sublevel 6 • d 10 • f 14

31. *No more than 2 electrons can occupy an orbital • *an electron will occupy the lowest possible sublevel

32. 9-The atomic number of the element : • 1- the elements are numbered consecutively from 1 to 106 coinciding with the number of protons in the nucleus • 2-H element number 1 has 1 proton in nucleus • 3- helium number 2 has 2 protons • Mg 12 protons • The atomic number of an element is the same as the number of protons in the nucleus ,the same as positive charge and also number of electrons in neutral atom.

33. 1-Hydrogen atom atomic number of the elements: • 1- The H atom consisting of a nucleus containing one proton and an electron • 2- orbital containing one electron ,is the simplest known atom • 3- The electron occupies an S orbital in the 1st energy level • 4-the electron doesn’t move in any definite path but rather in a random motion within its orbital forming an electron cloud about the • nucleus

34. 11- Isotpes of the elements • *atoms of an element having the same atomic number but different atomic masses are called • Isotopes of that element • *atoms of the isotopes of an element ,therefore have the same number of protons and elements but different numbers of neutrons

35. 12-atomic structure of the first twenty elements : • * the structure of the atoms of the 1st 20 elements ,arranged in the order of increasing atomic number (number of protons) • *the atoms of each succeeding element contain one more proton and one more electron than the atoms of the proceeding element. • *the number of neutrons in an atom also increases as we progress from the simpler elements to the more complex one • Periodic table page 23

36. Chemical bonding • 2- bonding and molecular structure • *chemical bond : • The attractive force that hold atoms together in compounds are called chemicalbonding

37. Bonding types

38. Bondingtypes • Ionic, covalent and metallic bonding • 1-ionic bond :term given to the electrostatic (charge-based)attractive forces which • Hold oppositely charged ions together • 2-Covalentbond : the sharing of electrons between two atoms that act s to hold the atoms together

39. *metallic bond : is found in metals .Atoms of the metal are bound to several neighbors, holding the atoms together but allowing electrons to move freely

40. Ionic bonding • *The ionic bond is the electrostatic force which attracts particles with opposite electrical charges • The formation of ions : • *Atoms can gain or lose electrons to become charged particles called ions • Cations: Are positively charged ions formed when an atom loses electrons

41. Anions: are negatively charged ions formed when an atom gain electrons • An ion is formed when an atom gain or losses one or more electron • M → M+ + e – • X + e- → x-

42. *If electron lost by M is gained by x ,the overall :reaction will be • M + X →M + + X - • M + + X - →M + X - • The ions attracted to each other because they have opposite charges ,the attraction is called an ionic bond or electrovalent bond .

43. Lewis structure : • *Lewis discover a Lewisstructure in which the chemical symbol for an atom is surrounded by a number of dots corresponding to the number of electrons in the valence shell of the atom . • e.gNa atom has one valence –shell-electron so its Lewis structure is • Na. • .

44. e.g. chlorine atom has 7 valence –shell electrons so its lewis structure is: • . • Cl: : • ..

45. The symbol = the nucleus plus all the inner shell electrons ,It called the core • E.g Al=13 • Electronic configration is 1S2 2S2 2P6 3S2 3P1 • . . • LewisstructureAl. • Here the valences are shown as a pair (the 2 3S electrons) And a single electron (the 3 P)

46. Octet rule : • * The octet rule is a statement of the stability • Of the nS2 –nP6 valence-shell configuration. • Atoms which can achieve this configuration by the addition of only a few electrons that is, tend to complete the octet . In adding electrons the atom becomes a negative ion. Thus the chloride ion is formed when one electron adds to a chlorine atom.

47. .. .. .Cl : + e - →[ :Cl : ] -.. .. • *Here the negative sign is written because the resulting particle is anion

48. In positive ions ,when has few valence electrons and has an octet in the second shell from the out side ,it tend to lose it • Valence electrons thereby exposing the octet. • In this way the resulting positive ion ends up with an octet in what is now its outer shell. • Thus the sodium ion tends to lose its valence • Electron to form a sodium ion : • Na(1S2 2S2 2P6 3S 1)→Na+ (1S2 2S2 2P6 )+e - • Na. → Na + + e -

49. Lewis structure and ionic compounds • * To write the Lewis structure for an ionic compound ,we write structures for the individual ions. Thus the Lewis structure for NaCl is: • .. • Na+ [ : Cl :] - • ..

50. :Note that • * The octet rule help us to predict stoichiometry that is , atomic combining ratio in ionic compounds.In the NaCl example one electron was transferred from one Na atom to • one Cl atom • .... • Na. + .Cl: → Na+ [ :Cl:] - • .. ..